Bonding theory, Lewis

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

Lewis s theory of the chemical bond was brilliant, but it was little more than guesswork inspired by insight. Lewis had no way of knowing why an electron pair was so important for the formation of covalent bonds. Valence-bond theory explained the importance of the electron pair in terms of spin-pairing but it could not explain the properties of some molecules. Molecular orbital theory, which is also based on quantum mechanics and was introduced in the late 1920s by Mul-liken and Hund, has proved to be the most successful theory of the chemical bond it overcomes all the deficiencies of Lewis s theory and is easier to use in calculations than valence-bond theory. 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

Lewis s theory also fails to account for the compound diborane, B2H6, a colorless gas that bursts into flame on contact with air. The problem is that diborane has only 12 valence electrons (three from each B atom, one from each H atom) but, for a Lewis structure, it needs at least seven bonds, and therefore 14 electrons, to bind the eight atoms together Diborane is an example of an electron-deficient compound, a compound with too few valence electrons to be assigned a valid Lewis structure. Valence-bond theory can account for the structures of electron-deficient compounds in terms of resonance, but the explanation is not straightforward. 

It was G. N. Lewis who extended the definitions of acids and bases still further, the underlying concept being derived from the electronic theory of valence. It provided a much broader definition of acids and bases than that provided by the Lowry-Bronsted concept, as it furnished explanations not in terms of ionic reactions but in terms of bond formation. According to this theory, an acid is any species that is capable of accepting a pair of electrons to establish a coordinate bond, whilst a base is any species capable of donating a pair of electrons to form such a coordinate bond. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. These definitions of acids and bases fit the Lowry-Bronsted and Arrhenius theories, and cover many other substances which could not be classified as acids or bases in terms of proton transfer. 

Clearly the explanation of the chemical bond given by Kossel cannot apply to homonuclear molecules such as CI2. Almost simultaneously with the publication of Kossel s theory, Lewis published a theory that could account for such molecules. Like Kossel, Lewis was impressed with the lack of reactivity of the noble gases. But he was also impressed by the observation that the vast majority of molecules have an even number of electrons, which led him to suggest that in molecules, electrons are usually present in pairs. In particular, he proposed that in a molecule such as CI2 the two atoms are held together by sharing a pair of electrons because in this way each atom can obtain a noble gas electron arrangement, as in the following examples ... 

There are several ways in which valence bond theory is superior to Lewis structures in... 

As predicted by elementary hybrid bonding theory, the multiple bonds of the chemist s Lewis-structure diagram are usually found to correspond to two distinct types of NBOs (1) sigma-type, having exact or approximate cylindrical symmetry about the bond axis (as discussed in Sections 3.2.5-3.2.7), and (2) pi-type, having a nodal mirror plane passing through the nuclei 44... 

Although the most naive form of valence-bond and Lewis-structure theory would not predict the paramagnetism of O2, the VB-like NBO donor-acceptor perspective allows us to develop an alternative localized picture of general wavefunctions, including those of MO type. Let us therefore seek to develop a general NBO-based configurational picture of homonuclear diatomics to complement the usual MO description. 

W. M. Latimer and W. H. Rodebush, J. Am. Chem. Soc. 42 (1920), 1419. This paper is often cited as the discovery of H-bonding (but see Jeffrey, note 5). Its title, Polarity and ionization from the standpoint of the Lewis theory of valence, reflects the strong influence of G. N. Lewis on all aspects of the early development of H-bond theory. 

This graduate-level text presents the first comprehensive overview of modern chemical valency and bonding theory, written by internationally recognized experts in the held. The authors build on the foundation of Lewis- and Pauling-like localized structural and hybridization concepts to present a book that is directly based on current ab initio computational technology. 

Nevertheless, even today, we often discuss the bonding of organic compounds in terms of Lewis structures and valence bond theory. 

At about the same time that Bronsted proposed his acid-base theory, Lewis put forth a broader theory, A base in the Lewis theory is the same as in the Brpnsted one, namely, a compound with an available pair of electrons, either unshared or in a tt orbital. A Lewis acid, however, is any species with a vacant orbital.1115 In a Lewis acid-base reaction the unshared pair of the base forms a covalent bond with the vacant orbital of the acid, as represented by the general equation... 

There is an implicit assumption contained in all of the above The two bonding electrons are of opposite spin. If two electrons are of parallel spin, no bonding occurs, but repulsion instead curve /, Fig. 5.1). This is a result of the Pauli exclusion principle. Because of the necessity for pairing in each bond formed, the valence bond theory is often referred to as the electron pair theory, and it forms a logical quantum-mechanical extension of Lewis s theory of electron pair formation. 

A major difference between the valence-bond theory of chemical bonding and molecular orbital theory is that the former assumes, like the Lewis approach, that the electrons in a bond are localized between the two... 

An advantage of VSEPR is its foundation upon Lewis electron-pair bond theory. No mention need be made of orbitals and overlap. If you can write down a Lewis structure for the molecule or polyatomic ion in question, with all valence electrons accounted for in bonding or nonbonding pairs, there should be no difficulty in arriving at the VSEPR prediction of its likely shape. Even when there may be some ambiguity as to the most appropriate Lewis structure, the VSEPR approach leads to the same result. For example, the molecule HIO, could be rendered, in terms of Lewis theory as ... 

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